MSE 170: Introduction to Materials Science and Engineering
Instructor

Prof. Christine Luscombe

Emailluscombe@u.washington.edu

Office302B Roberts Hall

Office hours

10:00am – 12:00pm, Monday, or by appointment

Course website

http://courses.washington.edu/mse170

Lab TAs:Peter Kazarinoff

peterkaz@u.washington.edu Mueller 168

Omid Veiseh

omid@u.washington.edu

Mueller 168

M 1:30-2:30pm

Steven Hau

skhau@u.washington.edu

Mueller 168

T 10:30-11:30am

1What is Materials Science and Engineering?

Aninterdisciplinary
fieldthat
addresses
thefundamental
relationships
betweentheProcessing, Structure and Propertiesof materials and develops them forthe desired technological application(Performance).

2

Materials through the ages
Stone Age (beginning of life – 3000 BC)

Feature: Using naturally occurring materials with only changes in shape3

Materials through the ages
Bronze Age (3000 BC – 1200 BC)
Copper and Tin alloy

Ability to modify materials byrefining (using heat), chemicalmodifications (alloying) andmechanical deformation (coldworking)
Imperfection (Ch. 4)
Diffusion (Ch. 5)
Phase diagram (Ch. 9)
Metal Processing (Ch. 11)4

Materials through the ages

Iron Age (1200 BC – Present)
Casting and alloying wasn’t perfected until
16th century
Mastery of Steel (Iron alloy) technology
enables Industrial Revolution in the 18thand 19th century
Ability to heat treat at high temperature, control
microstructure at different length scale and ability
to design specific microstructures for specificproperties
Phase transformation (Ch. 10)

5Materials through the ages
Plastic Age (1940 - Present)
Discovery of polymers, and the ability tosynthesize and process polymers.

6Materials through the ages
Silicon Age (1950 - Present)
Commercialization of silicon technology(integrated circuits, electronic devices, etc…)
leads to the information age, which gives
boost to human productivity

Electronic Prop. (Ch.18)
Thermal Prop. (Ch.19)
Magnetic Prop. (Ch.20)
Optical Prop. (Ch.21)

Ability to control alloying accurately,
ability to make thin films

7Future1. Nanotechnology- Synthesis and characterizations of nanomaterials and nanostructure

2. Biotechnology- biomimetics and biomaterials

3. Energy/Environmental- Next generation energy conversion

4. Information Technology- Materials informatics

8

What is Materials Science and Engineering?

Aninterdisciplinary
fieldthataddresses
thefundamentalrelationships
betweentheProcessing, Structure and Propertiesof materials and develops them forthe desired technological application(Performance).

9

Two Forms of Carbon

DiamondGraphite• A structure of carbon only
produced at high temperature andpressure.

• A structure of carbon that is inequilibrium (it is stable and will notchange form over time).

• The hardest known material.

• It is soft.

Atomic structure and interatomic bonding (Ch.2)
Crystallography (Ch. 3), Imperfection (Ch. 4)

10Ceramics
Ch. 12

Materials

MetalsCh. 11Polymers
Ch.14 and 15

Composite
Ch. 1611

Atomic structure
Nucleus
Neutron
Quarks GluonsElectronProtons
BOHR ATOM

orbital electrons:n = principal
quantum number12
n=3

Nucleus: Z = # protons= 1 for hydrogen to 94 for plutonium
Atomic mass A ≈ Z + N

N = # neutrons

12Electronic structure
Valence electrons determine all of the following properties:••••

ChemicalElectrical
Thermal
OpticalElectrons have wavelike and particulate properties.
• This means that electrons are in orbitals defined by a probability.
• Each orbital at discrete energy level determined by quantum numbers.

Quantum #Designationn = principal (energy level-shell)

K, L, M, N, O (1, 2, 3, etc.)

l = subsidiary (orbitals)s, p, d, f (0, 1, 2, 3,…, n -1)

ml = magnetic

1, 3, 5, 7 (-l to +l)

ms = spin

½, -½

13

Electronic structure
Number of electrons

Principal
quantum no.

Shelldesignation

Subshells

No. ofstates

Per subshell

Per shell

1Ks

1222

Ls

128p

36s

12p36d510s

12p36fdd510df

714ps

34M

N

18gf

32psdpspss
1234514

Electron energy states
Electrons...
• have discrete energy states• tend to occupy lowest available energy state.
4d4p

N-shell n = 4

3d4s

Energy

3p3s

2p2s
1s

M-shell n = 3

L-shell n = 2

K-shell n = 1
Adapted from Fig. 2.4,
15Callister 7e.

Electronic configuration
ex: Fe - atomic # = 26

1s2 2s2 2p6 3s2 3p6 3d 6 4s2

4d4p

valenceelectrons
N-shell n = 4

3d4s

Energy

3p3s

2p2s
1s

M-shell n = 3

L-shell n = 2K-shell n = 1
Adapted from Fig. 2.4,
16Callister 7e.

Survey of elements• Most elements: Electron configuration not stable.
Element

Atomic #Hydrogen
HeliumLithium
Beryllium
BoronCarbon...

123456Neon

10Sodium111s 2 2s 2 2p 6
1s 2 2s 2 2p 6 3s 1Magnesium

121s 2 2s 2 2p 6 3s 2Aluminum...

131s 2 2s 2 2p 6 3s 2 3p 1
...

Argon
...

18...

1s 2 2s 2 2p 6 3s 2 3p 6...

Krypton36

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6

Electron configuration
1s 1
1s 2(stable)
1s 2 2s 1
1s 2 2s 2
1s 2 2s 2 2p 1
1s 2 2s 2 2p 2...

• Why? Valence (outer) shell usually not filled completely.

(stable)(stable)(stable)Adapted from Table 2.2,Callister 7e.
17

The periodic table

Li BeNa Mg

inert gases

accept 1e

accept 2egive up 3eHgive up 2egive up 1e• Columns: Similar Valence Structure

HeO

F NeS

Cl Ar

K Ca Sc

Se Br KrRb SrTeYCs BaIXePo At Rn

Fr Ra

Electropositive elements:
Readily give up electrons
to become + ions.

Electronegative elements:
Readily acquire electrons
to become - ions.
Adapted from 18
Fig.
2.6, Callister 7e.

Electronegativity• Ranges from 0.7 to 4.0,
• Large values: tendency to acquire electrons.

Smaller electronegativity

Larger electronegativity

Adapted from Fig. 2.7, Callister 7e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical
Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.

19Bonding forces and energies• Attractive force, FArAttraction

Attractive force, FAE=" Fdrr

EN =0Repulsion

Force F+

• Repulsive force, FRr# F dr = # F dr + # F drN"!
Interatomic separation r

A"R"= EA + ER

Repulsive force, FR-r!

Net force, FN
Repulsive energy ER!

Interatomic separation r
Net energy
EN
Attractive energy EA

20

Properties from bonding• Bond length, r• Melting Temperature, Tm
Energyrr• Bond energy, EoEnergyror
smaller Tm

unstretched lengthoEolarger Tmr=Tm is larger if Eo is larger.

“bond energy”

21Properties from bonding: thermal expansion coefficient• Coefficient of thermal expansion, αlength,coeff. thermal expansion

Lounheated, T1Δ LΔ LL oheated, T2=α(T2-T1)• α ~ symmetry at ro
Energyunstretched lengthroEoEorα is larger if Eo is smaller.

larger αsmaller α

22Properties from bonding: modulus EF = kxr

Attraction

0Repulsion

Force F+23

Types of bonding: ionic• Occurs between + and - ions.
• Requires electron transfer.
• Large difference in electronegativity required.
• Example: NaCl

Na (metal)

Cl (nonmetal)

unstable

unstableelectron

Na (cation)
stable-+Coulombic

Cl (anion)
stableAttraction

24

Examples of ionic bonding• Predominant bonding in Ceramics
NaCl
MgO
CaF 2
CsCl

Give up electrons

Acquire electrons

Adapted from Fig. 2.7, Callister 7e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical
Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.

25

Covalent bonding•similar electronegativity ∴ share electrons•bonds determined by valence – s & p orbitals dominate bonding•Example: CH4
C: has 4 valence e-,needs 4 more

H: has 1 valence e-,needs 1 more

CH4H

H

C

Electronegativities
are comparable.

shared electrons
from carbon atomH
shared electrons

Hfrom hydrogenatoms

Adapted from Fig. 2.10, Callister 7e.

26

Metallic bonding•Ions in a sea of electrons•Attraction between free electrons and metal ions27

Ionic-covalent mixed bonding#(X A "X B )2 &"%( x ( 100 %)% ionic character = 1" e4%(%($'where XA & XB are Pauling electronegativities!
Example: MgO

XMg = 1.3
XO = 3.5&(3.5 '1.3 )2 #'$!
4% ionic character = $1 ' e! x (100%) = 70.2% ionic$!
%"

28

Secondary bonding
Arises from interaction between dipoles• Fluctuating dipoles

example: liquid H2

asymmetric electronclouds+-H2+-HH2HH

secondarybondingHsecondarybonding

Adapted from Fig. 2.13, Callister 7e.

• Permanent dipoles-molecule induced-general case:+-Cl

secondarybonding+secondarybonding

H-example: liquid HCl

H-example: polymer

secondarybonding

Adapted from Fig. 2.14,
Callister 7e.

Clsecondary bonding
29

SummaryType

Bond Energy

Comments

IonicLarge!

Non-directional (ceramics)

Covalent

Variable
Diamond (large)
Bismuth (small)

Directional(semiconductors, ceramics,polymer chains)

Metallic

Variable
Tungsten (large)
Mercury (small)Non-directional (metals)

Smallest

Directional
Interchain (polymer)
Intermolecular

Secondary

Ceramics(Ionic & covalent bonding)

Large bond energy
Large Tm and E, small αMetals(Metallic bonding)

Variable bond energy
Moderate Tm, E, and α

Polymers(Covalent & secondary)Directional properties,Secondary bonding
dominates
Small Tm and E, large α30

Energy and packing• Non dense, random packing

Energytypical neighborbond length
typical neighborbond energyrEnergy• Dense, ordered packing

typical neighborbond length

typical neighborbond energyrDense, ordered packed structures tend to havelower energies.
31Materials and packing
Crystalline materials...
• atoms pack in periodic, 3D arrays• typical of:-metals-many ceramics-some polymers
crystalline SiO2
Adapted from Fig. 3.22(a),Callister 7e.

SiOxygenNoncrystalline materials...
• atoms have no periodic packing• occurs for:-complex structures-rapid cooling"Amorphous" = Noncrystalline

noncrystalline SiO2
Adapted from Fig. 3.22(b),Callister 7e.

32

Crystals
Unit cell:smallest repetitive volume which contains the complete lattice pattern of acrystal.
7 crystal systems
14 crystal lattices

Lattice points: 3D array of points whichcoincides with atom positions.

Fig. 3.4, Callister 7e.

a, b, and c are the lattice constants
33

Simple cubic structure (SC)• Rare due to low packing density (only Po has this structure)• Close-packed directions are cube edges.
• Coordination # = 6(# nearest neighbors)

1 atoms/unit cell: 8 corners x 1/8(Courtesy P.M. Anderson)

34

Face centered cubic structure (FCC)• Atoms touch each other along face diagonals.
--Note: All atoms are identical; the face-centered atoms are shaded
differently only for ease of viewing.

ex: Al, Cu, Au, Pb, Ni, Pt, Ag• Coordination # = 12

Adapted from Fig. 3.1, Callister 7e.

4 atoms/unit cell: 6 face x 1/2 + 8 corners x 1/8(Courtesy P.M. Anderson)

35

Atomic packing factor (APF): FCC
Volume of atoms in unit cell*
APF =Volume of unit cell*assume hard spheres

volume2aatoms
unit cellaAdapted fromFig. 3.1(a),Callister 7e.

close-packed directions
contains 6 x 1/2 + 8 x 1/8 =4 atoms/unit cell

44 π R
3atom

3

APF =a3

volumeunit cell• APF for a FCC structure = 0.74
Maximum achievable APF36Theoretical density, ρDensity = ρ = Mass of atoms in unit cell
Total volume of unit cell
wherenA
VC NAρ =n = number of atoms/unit cell
A = atomic weight
VC = Volume of unit cell = a3 for cubic
NA = Avogadro’s number= 6.023 x 1023 atoms/mol

Ex: Cu (FCC)
A = 63.54 g/mol
R = 0.128 nm
n=4g

atomsunit cell
ρ=volumeunit cell

4 63.54a3

ρtheoretical = 8.9 g/cm3mol

ρactual = 8.94 g/cm3

6.023 x 1023

a = 2R 2atomsmol

37

Body centered cubic structure (BCC)• Atoms touch each other along cube diagonals.
--Note: All atoms are identical; the center atom is shaded
differently only for ease of viewing.

ex: Cr, W, Fe (α), Tantalum, Molybdenum

Adapted from Fig. 3.2,Callister 7e.

2 atoms/unit cell: 1 center + 8 corners x 1/8• Coordination # = 8(Courtesy P.M. Anderson)

38

Atomic packing factor (APF): BCC

volumeRatomsaunit cell

24 π (√ 3 a/4)33

atom

APF =a3

volumeunit cell• APF for a FCC structure = 0.68
39

FCC stacking sequence• ABCABC... Stacking Sequence• 2D Projection

BA sites
B sites

ABCB

CBB
CB

BC sites• FCC Unit Cell

ABC40

Hexagonal close-packed structure (hcp)• ABAB... Stacking Sequence• 3D Projection• 2D Projection

ca

A sitesToplayer

B sites

Middle layer

A sites

Bottom layer

Adapted from Fig. 3.3(a),Callister 7e.

• Coordination # = 12
• APF = 0.74

6 atoms/unit cell
ex: Cd, Mg, Ti, Zn• c/a = 1.633
41